Transcript Ch 7 ppt - mvhs

Chemistry, The Central Science, 11th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 7
Periodic Properties
of the Elements
John D. Bookstaver
St. Charles Community College
Cottleville, MO
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Periodic Trends
• In this chapter, we will rationalize observed
trends in
– Sizes of atoms and ions
– Ionization energy
– Electron affinity
– Electronegativity
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Periodic Trends Key Words
• Principal Energy Levels: The more the number
of principal energy levels, the bigger the size
of atoms.
• Nuclear Charge (# of p in an atom): Results into
increased attraction on electrons. Causes atomic
radius to decrease.
• Shielding Effect: Electrons present between
nucleus and outermost energy level (all
electrons except for valence electrons).
Periodic Trends Key Words
Shielding electrons tend to increase atomic size
by reducing the attractive force on outermost
electrons.
• Effective Nuclear Charge: Force of attraction
felt by the outermost (valence e) from the
protons in the nucleus. Effective nuclear
charge depends upon the two counteractive
factors of nuclear charge and shielding effect.
A high effective nuclear charge means smaller
ionic radius (greater attraction on the
outermost electrons).
Effective Nuclear Charge
• In a many-electron atom,
electrons are both
attracted to the nucleus
and repelled by other
electrons.
• The nuclear charge that
an electron experiences
depends on both factors.
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Effective Nuclear Charge
The effective nuclear
charge, Zeff, is found this
way:
Zeff = Z − S
where Z is the atomic
number and S is a
screening constant,
usually close to the
number of inner
electrons.
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Effective Nuclear Charge and Shielding
What Is the Size of an Atom?
The bonding atomic
radius is defined as
one-half of the
distance between
covalently bonded
nuclei.
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Sizes of Atoms
Bonding atomic radius
tends to…
…decrease from left to
right across a row
(due to increasing Zeff).
…increase from top to
bottom of a column
(due to increasing value of
n).
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/atomic4.swf
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Periodic Trend for Atomic Radius
Going down a group- ENC is lower; size is larger:
Higher energy levels have larger orbitals
Shielding - core e- block the attraction between the nucleus
and the valence eAcross a group- ENC is higer; size is smaller:
Increased nuclear charge without additional shielding
pulls e- in tighter
Atomic Radius (pm)
250
200
150
100
50
0
0
5
10
Atomic Numbe r
15
20
Atomic Radius
• Why larger going down?
– Higher energy levels have larger orbitals
– Shielding - core e- block the attraction between the
nucleus and the valence e-
• Why smaller to the right?
– Increased nuclear charge without additional shielding
pulls e- in tighter
Explain to your shoulder partner the
atomic radius trend:
a. Across a period
b. Down a group
Be sure to use the following key words:
• N (number of energy levels)
• Nuclear Charge
• Shielding Effect
• Effective Nuclear Charge
Sizes of Ions
• Ionic size depends
upon:
– The nuclear charge.
– The number of
electrons.
– The orbitals in which
electrons reside.
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Sizes of Ions
• Cations are smaller
than their parent
atoms.
– The outermost
electron is removed
and repulsions
between electrons
are reduced.
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Sizes of Ions
• Anions are larger
than their parent
atoms.
– Electrons are added
and repulsions
between electrons
are increased.
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Sizes of Ions
• Ions increase in size as
you go down a
column.
– This is due to
increasing value of n.
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Sizes of Ions
• In an isoelectronic series, ions have the same
number of electrons.
• Ionic size decreases with an increasing nuclear
charge.
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Find a new partner (someone who you have never worked with before!)
and do the following pair-share activity:
Explain to your partner periodic trend for size of ions
Your partner explains to you the group trend for size of ions
Make sure to use the following key words:
•N (number of energy levels)
•Nuclear charge
•Shielding effect
•Effective Nuclear Charge
Now find another group and explain your partner’s reasoning for the trend
to them, while your partner explains your reasoning for the trend to them.
Ionization Energy
• The ionization energy is the amount of
energy required to remove an electron
from the ground state of a gaseous atom or
ion.
• A(g) A+ + e
– The first ionization energy is that energy
required to remove first electron.
– The second ionization energy is that energy
required to remove second electron, etc.
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Ionization Energy
• It requires more energy to remove each successive
electron.
• When all valence electrons have been removed, the
ionization energy takes a quantum leap. Why? Hint:
where are the electrons being removed from?
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Trends in First Ionization Energies
• As one goes down a
column, less energy is
required to remove the
first electron.
– For atoms in the same
group, Zeff is essentially
the same, but the
valence electrons are
farther from the nucleus.
http://nuweb.neu.edu/bmaheswaran/phyu121/data/ch09/anim/anim0903.htm
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Trends in First Ionization Energies
• Generally, as one goes
across a row, it gets
harder to remove an
electron.
– As you go from left to
right, Zeff increases.
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Trends in First Ionization Energies
However, there are
two apparent
discontinuities in this
trend.
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Trends in First Ionization Energies
• The first discontinuity
occurs between Groups
2 and 13.
• In group 13, the electron
is removed from a porbital rather than an sorbital.
– The electron removed is
farther from nucleus.
– There is also a small
amount of repulsion by
the s electrons.
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Pair-Share Activity
With your elbow partner, discuss the following
question. You might be called upon to share
your explanation with the whole class:
• Which element has a higher Ionization energyBe or B? Why?
Trends in First Ionization Energies
• The second anomaly
occurs between Groups
5 and 6. Following the
trend, the IE of group 6
elements should be
higher than group 5, but
IE of group 6 elements is
lower than that of group
5.
• The electron removed comes
from doubly occupied orbital.
• Repulsion from the other
electron in the orbital aids in
its removal.
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Sample Problem
• Write the answer
to the following
sample problem on
a piece of paper.
You will be grading
your elbow
partner’s paper at
the end.
• Question: Which
element has a
higher Ionization
energy- N or O?
Why?
Electron Affinity
Electron affinity is the energy change
accompanying the addition of an electron to a
gaseous atom:
Cl (g) + e−  Cl−
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Trends in Electron Affinity
In general, electron
affinity becomes more
exothermic (larger –
value) as you go from
left to right across a
row.
http://www.youtube.com/watch?v=bPB0xThmpkg&feature=related
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Trends in Electron Affinity
There are again,
however, two
discontinuities
in this trend.
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Trends in Electron Affinity
• The first occurs
between Groups 1 and
2, where EA of group 2
is lesser than group 1.
– The electron is farther
from nucleus and feels
repulsion from the selectrons.
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Trends in Electron Affinity
• The second occurs
between Groups 14and
15.
– Group 15 has no empty
orbitals (only half filled
p orbitals).
– The extra electron must
go into an already
occupied orbital,
creating repulsion.
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Sample Problem
• Why is the general • Which of the
periodic trend for
following elements
EA? What is the
has a higher EA?
group trend for EA? • Na or Mg?
Why?
• P or Si?
• Why?
Electronegativity
•Increases from L to R across a period and decreases down
a group.
•Electronegativity is defined as tendency to attract electrons
but it is different from electron affinity in that electro negativity
is used in context of an element BONDED IN A COVALENT
COMPOUND, while electron affinity is generally attributed to
an atom by itself.
•Another difference is that electro negativity is a measure of
affinity for electrons in debye scale, while electron affinity is
the actual amount of energy released.
http://www.youtube.com/watch?v=93G_FqpGFGY
http://www.youtube.com/watch?v=WyfwRPBw62s
Good Link on Periodic Trends
Properties of Metal, Nonmetals,
and Metalloids
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Metals versus Nonmetals
Differences between metals and nonmetals tend
to revolve around these properties.
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Metals versus Nonmetals
• Metals tend to form cations.
• Nonmetals tend to form anions.
• Metallic character increases down a group and
decreases across a period.
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Metals
• Compounds formed
between metals and
nonmetals tend to be
ionic.
• Metal oxides tend to
be basic.
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Nonmetals
• These are dull, brittle
substances that are
poor conductors of heat
and electricity.
• They tend to gain
electrons in reactions
with metals to acquire a
noble gas configuration.
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Nonmetals
• Substances containing
only nonmetals are
molecular
compounds.
• Most nonmetal oxides
are acidic.
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Metalloids
• These have some
characteristics of
metals and some of
nonmetals.
• For instance, silicon
looks shiny, but is
brittle and fairly poor
conductor.
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Group Trends
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Alkali Metals
• Alkali metals are soft,
metallic solids.
• The name comes from
the Arabic word for
ashes.
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Alkali Metals
• They are found only in compounds in nature, not
in their elemental forms.
• They have low densities and melting points.
• They also have low ionization energies.
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Alkali Metals
Their reactions with water are famously exothermic.
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Alkali Metals
• Alkali metals (except Li) react with oxygen to form
peroxides.
• K, Rb, and Cs also form superoxides:
K + O2  KO2
• They produce bright colors when placed in a
flame.
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Alkaline Earth Metals
• Alkaline earth metals have higher densities and
melting points than alkali metals.
• Their ionization energies are low, but not as low
as those of alkali metals.
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Alkaline Earth Metals
• Beryllium does not
react with water and
magnesium reacts
only with steam, but
the others react
readily with water.
• Reactivity tends to
increase as you go
down the group.
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Group 6A
• Oxygen, sulfur, and selenium are nonmetals.
• Tellurium is a metalloid.
• The radioactive polonium is a metal.
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Oxygen
• There are two allotropes of
oxygen:
– O2
– O3, ozone
• There can be three anions:
– O2−, oxide
– O22−, peroxide
– O21−, superoxide
• It tends to take electrons from
other elements (oxidation).
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Sulfur
• Sulfur is a weaker
oxidizer than oxygen.
• The most stable
allotrope is S8, a
ringed molecule.
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Group VIIA: Halogens
• The halogens are prototypical nonmetals.
• The name comes from the Greek words halos and
gennao: “salt formers”.
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Group VIIA: Halogens
• They have large, negative
electron affinities.
– Therefore, they tend to oxidize
other elements easily.
• They react directly with
metals to form metal halides.
• Chlorine is added to water
supplies to serve as a
disinfectant
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Group VIIIA: Noble Gases
• The noble gases have astronomical ionization
energies.
• Their electron affinities are positive.
– Therefore, they are relatively unreactive.
• They are found as monatomic gases.
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Group VIIIA: Noble Gases
• Xe forms three
compounds:
– XeF2
– XeF4 (at right)
– XeF6
• Kr forms only one stable
compound:
– KrF2
• The unstable HArF was
synthesized in 2000.
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